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Bonding and Molecular Structure: Valence e- and Bonding Covalent Ionic Bond Energy & Length

Chapter 6 Chemical Bonding. Bonding and Molecular Structure: Valence e- and Bonding Covalent Ionic Bond Energy & Length Structure, Shape & Polarity of Compounds. What is a Bond?. A force that holds atoms together. Why? We will look at it in terms of energy.

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Bonding and Molecular Structure: Valence e- and Bonding Covalent Ionic Bond Energy & Length

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  1. Chapter 6 Chemical Bonding Bonding and Molecular Structure: Valence e- and Bonding Covalent Ionic Bond Energy & Length Structure, Shape & Polarity of Compounds

  2. What is a Bond? • A force that holds atoms together. • Why? • We will look at it in terms of energy. • Bond energy the energy required to break a bond. • Why are compounds formed? • Because it gives the system the lowest energy.

  3. Covalent compounds? • The electrons in each atom are attracted to the nucleus of the other. • The electrons repel each other, • The nuclei repel each other. • The reach a distance with the lowest possible energy. • The distance between is the bond length.

  4. Thus Hydrogen is Diatomic! Bond Formation

  5. e- Covalent Character

  6. He2 . . E He + He . . Inter-nuclear Distance Why Isn’t Helium Diatomic?

  7. F + F F2 2p____ ____ ___ ___ ____ ____ 2p2s ____ ____ 2s F F

  8. Ionic Bonding • An atom with a low ionization energy reacts with an atom with high electron affinity. • The electron moves. • Opposite charges hold the atoms together.

  9. Li + Cl1s22s1 [Ne] 3s23p52s ___ 3p _____ _____ ___1s _____ 3s _____[Ne]

  10. Li + Cl 2s ___ 3P _____ _____ _____1s _____ 3s _____[Ne]

  11. LiCl2s ___3P _____ _____ _____1s _____ 3s _____ [Ne]

  12. Electronegativity The difference between ionic and covalent bonds. Describes the relative ability of an atom within a molecule to attract a shared pair of electrons to itself.

  13. Electronegativity Pauling electronegativity values, which are unit-less, are the norm.

  14. ElectronegativityRange from 0.7 to 4.0

  15. Bond: A - B DEN = | ENA - ENB |

  16. Bond Character “Ionic Bond” - Principally Ionic Character “Covalent Bond” - Principally Covalent Character

  17. covalent ionic EN ~0 1.7 ~4 Determining Principal Character of Bond

  18. F - F EN = 0 Non-polar

  19. N - O EN = |3.0 - 3.5| = 0.5 O N Slightly polar

  20. Ca - O  EN = |1.0 - 3.5| = 2.5 Ca O Ionic Bond with somecovalent character

  21. Electronegativity • D is known for almost every element • Gives us relative electronegativities of all elements. • Tends to increase left to right. • decreases as you go down a group. • Noble gases aren’t discussed. • Difference in electronegativity between atoms tells us how polar.

  22. Polar Covalent Ionic Electronegativity difference Bond Type Zero Covalent Covalent Character decreases Ionic Character increases Intermediate Large

  23. Dipole Moments • A molecule with a center of negative charge and a center of positive charge is dipolar (two poles), • or has a dipole moment. • Center of charge doesn’t have to be on an atom. • Will line up in the presence of an electric field.

  24. d+ d- H - F How It is drawn

  25. Which Molecules Have Them? • Any two atom molecule with a polar bond. • With three or more atoms there are two considerations. • There must be a polar bond. • Geometry can’t cancel it out.

  26. Ionic Radii -- Cations

  27. Ionic Radii -- Anions

  28. Molecular Polarity MgBr2 Mg - Br EN = |1.2 - 2.8| = 1.6 Mg Br Br Covalent BOND w/much ionic character, BUT NON-POLAR molecule

  29. Lewis Structures

  30. The most important requirement for the formation of a stable compound is that the atoms achieve noble gas e- configuration

  31. Valence Shell ElectronPair Repulsion Model(VSEPR) The structure around a given atom is determined principally by minimizing electron-pair repulsions

  32. Electron pairs Bond Angles Underlying Shape 2 180° Linear 3 120° Trigonal Planar 4 109.5° Tetrahedral 90° & 120° Trigonal Bipyramidal 5 6 90° Octagonal VSEPR

  33. LEWIS STRUCTURES • : draw skeleton of species • : count e- in species • : subtract 2 e- for each bond in skeleton • : distribute remaining e-

  34. Distinguish Between ELECTRONIC Geometry & MOLECULAR Geometry

  35. CH4 Bond angle = 109.50 Electronic geometry: tetrahedral Molecular geometry: tetrahedral

  36. H3O+ Bond angle ~ 1070 Electronic geometry: tetrahedral Molecular geometry: trigonal pyramidal

  37. H2O Bond angle ~ 104.50 Electronic geometry: tetrahedral Molecular geometry: bent

  38. NH2- Bond angle ~ 104.50 Electronic geometry: tetrahedral Molecular geometry: bent

  39. “Octet Rule” holds for connecting atoms, but may not for the central atom.

  40. BaI2 Bond angle =1800 Electronic geometry: linear Molecular geometry: linear

  41. BF3 Bond angle =1200 Electronic geometry: trigonal planar Molecular geometry: trigonal planar

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